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33 Cards in this Set
- Front
- Back
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Democritus |
- Greek philosopher that theorised that matter is composed of indivisible particles. However, no concrete scientific evidence was given |
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Dalton's model of the atom (1808) |
- All matter consists of tiny indivisible particles, called atoms - Atoms cannot be created or destroyed - An element consists of atoms of 1 type only - Atoms combine (in fixed ratios) to form compounds |
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Thomson (1897) |
- Cathode ray tube experiment showed existence of negatively charged particles 2000 times smaller than the smallest atom H: the electron - Thus he showed that atoms are destructible (not the smallest particle) - Plum pudding model: Electrons, scattered in a positively charged sponge-like substance |
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Rutherford (1910) |
- He shot alpha particles (=He nuclei): emitted by large atoms with too many protons to be stable: through a gold foil - Large number of undeflected particles suggests that the atom is mainly empty space - Deflection of some particles suggests that there is a positively charged center, the nucleus, which is very small (few alpha particles bounce back) |
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Bohr (1913) |
- Niels Bohr examined the line emission spectrum of the Hydrogen atom and proposed an atomic model of the Hydrogen atom, resembling the solar system: the negative electron moves in an orbit (energy level) around the positive nucleus. The electrostatic force of attraction prevents the e- from leaving the atom |
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Sub-atomic particles |
- Electrons - Protons - Neutrons |
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Nucleons |
- Protons - Neutrons |
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Atomic Mass Unit |
- 1 12th of the mass of a C-12 atom in the ground state - Represents the mass of 1 proton or neutron |
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The Nuclear Symbol Notation |
A - Z = Neutrons |
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Isotopes |
- Atoms of the same element with the same # of protons and different # of neutrons (different mass number) |
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Chemical Properties of an Isotope |
- Isotopes of an element have the same number of electrons and, as it is the number of electrons that determines the chemical properties of a substance, all isotopes of the same element will have the same chemical properties |
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Physical Properties of an Isotope |
- Physical properties often also depend on the mass of the particles and so different isotopes of an element will have different physical properties, for example mass, density, boiling and melting point |
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Applications of Isotopes |
- Nuclear Medicine: Radioisotopes are used in nuclear medicine for diagnosis (tracers) and treatment a.) Info about the functioning of a person's specific organs b.) Treat disease PET (Position Emission Tomography) scanners give 3D imaging of tracer concentration in the body, by detecting y-rays emitted by the tracers, thus producing "slice" images and detecting i.e. cancers - Carbon Dating: Carbon radioisotopes are used in cosmic, geological and archaeological dating. |
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Ions |
- Atoms that have a positive or negative charge, after losing or gaining electrons - Ions for when an atom gains or loses electron(s) in order to achieve the same stable electronic configuration as the atoms of the nearest noble gas - Atoms of the metallic elements tend to form ions by losing electrons a.) These positive charged ions are called cations - Atoms of non-metallic elements tend to form ions by gaining electrons b.) These negative charged ions are called anions - Chemical differences between elements and their ions e.g. Na is explosive in water, Na cations are not |
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Relative Atomic Mass |
- Determined by the mass spectrometer - To calculate Ar, relative abundances of all isotopes must be taken into consideration |
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Mass spectrometer |
- Can measure the mass of individual atoms in a sample. There can be isotopes of an element in any naturally occurring sample. |
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The electromagnetic spectrum (EM spectrum) |
- Atoms of different elements give out -upon heating- light of distinctive colour - Electromagnetic radiation: Radiation that has properties of both an electric and magnetic wave and that travels through a vacuum with the speed of light, c. - Properties of waves a.) Travel through a vacuum at the speed of light b.) Wavelength: Distance between two successive crests c.) Frequency: The number of waves which pass a particular point in one second |
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Range of the Electromagnetic Spectrum |
- Radio > Microwave > Infrared > Visible > Ultraviolet > X-ray > Gamma Ray - Radio = high wavelength low velocity - Gamma ray = Low wavelength high velocity |
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Forms of the Electromagnetic Spectrum |
- Continuous spectrum: Contains all colours / wavelengths / λ / v and energy levels. When sunlight passes through a prism it produces a continuous spectrum, much like a rainbow - Line spectrum: Elements will emit light in discrete λ / f / colours thus producing a line spectrum when excited by the passage of an electric discharge through their gas or vapour. For many metals the same effect (discrete lines) can be observed when their compounds are heated directly in a Bunsen burner. |
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Atomic Absorption and Emission Line Spectra |
- Atoms first absorb energy (absorption spectrum), and then as they lose it they emit light - This light, viewed through a spectroscope, will give the line emission spectrum of that element (discrete lines!), which is unique to that particular element - like a fingerprint or barcode. |
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Bohr's Model and the Hydrogen Emission Spectrum |
- Electrons (e-) move in circular orbits around the nucleus, with fixed amounts of potential energy - These e- are usually found in their ground state - In order to move to an orbit further away from the nucleus, the e- must absorb a fixed quantity of energy known as a quantum or photon to do work against the attraction of the positively charged nucleus. The electron is now said to be in an excited state - The excited state is unstable and the e- soon falls back to the ground state, emitting a photon, which is a specific amount of energy, with a specific v and λ - The emission spectrum is formed when e- return to a lower energy level emitting light of specific frequencies - If the incoming energy is as high as an atom's ionisation energy, it will lose the electron |
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Electron Transitions and Energy of Photons Emitted |
1.) Transitions to the 1st Energy Level correspond to an increased energy change and are found in the UV region of the EM spectrum 2.) Hydrogen produces visible light when electrons fall down to the 2nd energy level 3.) Transitions to 3rd energy levels correspond to IR radiation |
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Atomic Orbitals |
- 3D regions around the nucleus in which there is a 90% probability of the electron to be found |
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1st energy level |
- n = 1 - The 1st energy level consits of one s orbital - Shape: sphere - One s orbital can hold max 2 e- |
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2nd energy level |
- n=2 - Is split into 2 sub-levels: s, p - two s sub-level has 1s orbital - two s orbitals larger than 1s. Shape: spherical - three p orbitals (px, py, pz) have a dumbbell shape, and are arranged at perpendicular angles to each other, along the x, y, and z axes. They are called degenerate because they are of equal E - All orbitals can hold max 2 electrons |
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3rd energy level |
- n=3 - Split into three sub-levels s, p, d |
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Electron configuration |
- The periodic table can be subdivided into four blocks (s, p, d, and f) corresponding to the outer electron configurations of the elements in these blocks - 4s orbital is removed before 3d orbital |
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Orbital Diagrams |
- Block Diagrams ([↑↓]) |
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Full Electron Configuration Diagram |
- Written out (1s2, 2s2, 2p6, 3s3, 3p6, 4s2, 3d6) - Condensed = ([Ar] 4s2, 3d6) |
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Exceptions to electron configuration |
- Cr, Cu - Half filled or fully filled shells are more stable - In case of d shell, since the s electron of the next period is close in energy, it is possible to promote an s orbital electron to the d shell of the previous period, thus filling or half-filling the d shell, minimising repulsions with singly occupied atomic orbitals. This filled or half-filled state is more energetically stable. |
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First Ionisation Energy |
- The minimum amount of energy required to remove one electron from a gaseous atom - IE increases across a period - IE decreases down a group |
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Evidence for sub-levels |
- Drop in IE from group 2 to group 3, as it is energetically easier to remove an electron from the p sub-level than the s e- of the same energy level a.) electrons in p orbitals are further away from the nucleus and have lower energy than s electrons of the same energy level, so they are easier to remove - Drop in IE group 15 to group 16, as it is energetically easier to remove a paired electron in group 16 rhan an unpaired electron in group 15 a.) an electron in a doubly occupied orbital is repelled by its paire e- and so is easier to remove than an e- in a half-filled orbital |
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Successive Ionisation Energies |
- As electrons are removed closer to the nucleus, they require more energy to remove - Successive ionisation energies increase, with large "jumps" from different energy levels |
